Lewis structure & molecular geomtery

Lewis Structure

• A Lewis structure depicts the distribution of valence electrons in a molecule.
• Steps:
  • Count total valence electrons.
  • Connect atoms with single bonds.
  • Subtract electrons used for bonds.
  • Complete octets for exterior atoms (8 electrons).
  • Subtract electrons used.
  • Place remaining electrons on the central atom.
  • If central atom lacks octet, try using multiple bonds.

Lewis Structure Examples

• CH₄, NH₃, H₂O all have 8 valence electrons.

Resonance Structures

• Resonance refers to the delocalization of electrons in a molecule.
• Molecules with resonance structures exhibit multiple Lewis structures, each contributing to the overall structure.
• Example: CO₂ has three resonance structures.
• The actual structure is a blend (resonance hybrid) of all resonance forms.

Formal Charge

• Formal charge represents the hypothetical charge an atom would have in a molecule if electrons were equally shared.
• Calculation: Formal Charge = Group Number - (Dashes + Dots).
• The sum of all formal charges in a molecule should equal the overall charge.

Resonance Stability

• Resonance forms with the most zero formal charges are more stable.
• Negative charges are preferred on more electronegative atoms.
• Resonance forms with fewer non-zero formal charges are more stable and contribute more to the resonance hybrid.

VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular geometry based on electron pair repulsion.
• Key Points:
  • Electron pairs repel each other to maximize distances, resulting in specific molecular shapes.
  • A correct Lewis structure is required to apply VSEPR.
  • Regions of Electron Density: Each single bond, lone pair, double bond, or triple bond represents one region of electron density.

Electron Pair Geometry

• Two electron pairs: Linear geometry (180°).
• Three electron pairs: Trigonal planar geometry (120°).
• Four electron pairs: Tetrahedral geometry (109.5°).
• Five electron pairs: Trigonal bipyramidal geometry (90° and 120°).
• Six electron pairs: Octahedral geometry (90°).

Molecular Geometry

• Molecular geometry refers to the arrangement of atoms in a molecule.
• To determine molecular geometry, first identify the electron pair geometry and then ignore lone pairs to see the arrangement of atoms.

Nonpolar Compounds

• Linear AB₂, trigonal planar AB₃, square planar and tetrahedral AB₄, trigonal bipyramidal AB₅, and octahedral AB₆ are generally nonpolar.
• Any molecule not fitting these descriptions is likely polar.

Lewis Structures

• A Lewis structure shows the distribution of valence electrons in a molecule or ion.
• Follow these steps to draw a Lewis structure:
  • Count the total number of valence electrons.
  • Join atoms symmetrically with single bonds.
  • Subtract electrons used in bonds.
  • Complete the octet (8 electrons) on exterior atoms.
  • Subtract used electrons.
  • Place remaining electrons on the central atom.
  • If the central atom lacks an octet, consider using multiple bonds.
• Examples: CH4, NH3, H2O (all have 8 valence electrons)

Resonance Structures

• Consider the CO2 molecule:
  • There are multiple possible Lewis structures due to the double bonds.
  • The actual structure is a blend of all resonance forms, called a resonance hybrid.
• The stability of each resonance form is determined by formal charge.

Formal Charge

• Formal charge is the charge an atom would have if bonding electrons were equally shared.
• Formula: Formal Charge = Group Number - (Dashes + Dots)
  • Dashes represent bonds, dots represent lone pairs.
• The sum of all formal charges should equal the total charge of the molecule/ion.
• Resonance forms with more zero formal charges and negative charges on the most electronegative atoms are more stable.

VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) theory explains the shape of molecules based on electron pair repulsion.
• Electron pairs (bonds and lone pairs) want to maximize distance from each other.
• Steps:
  • Draw the Lewis structure.
  • Determine the number of electron density regions around the central atom (bonds or lone pairs).
  • This gives the electron pair (or basic) geometry.
  • Ignore lone pairs to get the molecular geometry.

Electron Pair Geometry

• 2 electron pairs: Linear (180°)
• 3 electron pairs: Trigonal planar (120°)
• 4 electron pairs: Tetrahedral (109.5°)
• 5 electron pairs: Trigonal bipyramidal (90° and 120°)
• 6 electron pairs: Octahedral (90°)

Molecular Geometry

• The shape of a molecule is determined by the arrangement of atoms, ignoring lone pairs.
• Examples:
  • Tetrahedral (CH4)
  • Trigonal pyramidal (NH3)
  • Bent (H2O)

Nonpolar Compounds

• Linear AB2 (BeCl2)
• Trigonal planar AB3 (BF3)
• Square planar and Tetrahedral AB4 (CH4, SiCl4)
• Trigonal bipyramidal AB5 (PCl5)
• Octahedral AB6 (SF6)
• Any other combination is polar.

Gilbert N. Lewis

• American physical chemist.
• Discovered the covalent bond and the concept of shared electron pairs.
• Known for Lewis dot structures and his contributions to valence bond theory.
• Coined the term "photon" in 1926.
• His research covered areas like chemical thermodynamics, photochemistry, and isotope separation.
• Died in his lab of cyanide poisoning on March 23, 1946.

Explore the fundamental concepts of Lewis structures and resonance in this chemistry quiz. Learn the steps to draw Lewis structures, understand resonance forms, and calculate formal charge. Test your knowledge on how these concepts apply to molecules like CH₄ and CO₂.